Electrochemistry and corrosion notes — Unit 3

Electrochemistry is the branch of chemistry that deals with the interconversion of electrical energy and chemical energy through redox (reduction-oxidation) reactions.

⇒Redox Reactions Fundamentals

• Oxidation: The LOSS of electrons. The oxidation state increases. (e.g., Zn → Zn^2+ + 2e^-)
• Reduction: The GAIN of electrons. The oxidation state decreases. (e.g., Cu^2+ + 2e^- → Cu)

⇒Types of Electrochemical Cells

The Daniel cell is the classic example of a Galvanic cell. It harnesses the spontaneous redox reaction between Zinc and Copper.

⇒Construction & Reactions

It consists of a Zinc rod dipped in ZnSO4 solution (Anode) and a Copper rod dipped in CuSO4 solution (Cathode), connected by a salt bridge.

// At Anode (Oxidation): Zinc dissolves into the solution
Zn(s) → Zn^2+(aq) + 2e^-

// At Cathode (Reduction): Copper ions plate onto the rod
Cu^2+(aq) + 2e^- → Cu(s)

// Overall Cell Reaction:
Zn(s) + Cu^2+(aq) → Zn^2+(aq) + Cu(s)

⇒The Salt Bridge

A U-tube containing a semi-solid paste of an inert electrolyte like KCl or NH4NO3 in agar-agar gel. Its functions are critical:

• It completes the internal electrical circuit by allowing ions to flow.
• It maintains electrical neutrality in both half-cells. Without it, excess Zn^2+ would build up at the anode and stop the reaction instantly.
• It prevents the two solutions from mixing mechanically.

⇒Electrode Potential (E)

When a metal is dipped into a solution of its own ions, a potential difference develops at the interface between the metal and the solution. This is the tendency of the electrode to either lose or gain electrons.

• Standard Electrode Potential (E°): Measured under standard conditions: 298 K temperature, 1 Molar concentration, and 1 atm pressure.

⇒Electromotive Force (EMF)

The EMF (or Cell Potential, E_cell) is the potential difference between the two electrodes of a galvanic cell. It is the driving force that pushes electrons through the external circuit.

E°_cell = E°_cathode - E°_anode

// Example for Daniel Cell:
E°_Cu = +0.34 V (Cathode)
E°_Zn = -0.76 V (Anode)
E°_cell = 0.34 - (-0.76) = 1.10 Volts

If E_cell is positive, the reaction is spontaneous and the cell will generate electricity. If negative, the cell will not work.

Standard potentials (E°) are only valid at exactly 1 Molar concentration. In reality, batteries discharge, and their concentrations change. Walther Nernst provided the equation to calculate the electrode potential at ANY concentration and temperature.

⇒Mathematical Formulation

For a general reduction reaction: M^n+ + ne^- → M

E = E° - (RT / nF) * ln( [Product] / [Reactant] )

// At standard 298K, substituting R, T, and F, and converting ln to log10:
E = E° - (0.0591 / n) * log10( 1 / [M^n+] )

Where:
E = Electrode potential under non-standard conditions
E° = Standard electrode potential
n = Number of electrons transferred in the reaction
F = Faraday's constant (96500 C/mol)
[M^n+] = Molar concentration of the metal ions

⇒EMF of a Complete Cell

The Nernst equation can be applied to the entire cell to calculate the exact voltage a battery will output as it discharges.

E_cell = E°_cell - (0.0591 / n) * log10( [Anode Ions] / [Cathode Ions] )

You cannot measure the absolute potential of a single half-cell. You can only measure the difference between two half-cells. Therefore, we assign an arbitrary potential of 0.00 V to a standard reference electrode, and measure everything else against it.

⇒Standard Hydrogen Electrode (SHE)

The primary reference electrode. It consists of a platinum wire coated with platinum black, dipped in a 1M HCl solution. Hydrogen gas at 1 atm is bubbled through it.

• Reaction: 2H+ + 2e- ⇌ H2(g)
• Potential: E° = 0.00 V (Arbitrarily assigned at all temperatures)
• Drawbacks: It is extremely difficult to maintain exactly 1 atm of H2 gas and 1M concentration continuously. It is fragile and easily poisoned.

⇒Secondary References: Calomel Electrode

Because SHE is impractical for daily lab use, secondary electrodes with known, highly stable potentials are used. The Saturated Calomel Electrode (SCE) is the most common.

It consists of Mercury covered by a paste of Mercurous Chloride (Calomel, Hg2Cl2), immersed in a saturated KCl solution. Its potential is fixed at +0.2422 V against SHE.

A commercial battery is simply one or more galvanic cells connected in series to provide a steady DC voltage for consumer use.

⇒Classification of Batteries

⇒The Dry Cell (Leclanché Cell)

The classic primary cell used in clocks and remotes.

• Anode: A Zinc container.
• Cathode: A Graphite rod surrounded by MnO2.
• Electrolyte: A paste of NH4Cl and ZnCl2.
• Voltage: ~1.5 V.

The Lead-Acid battery is the quintessential secondary battery, universally used in automotive engines to start cars. It provides high surge currents.

⇒Construction

• Anode: Spongy metallic Lead (Pb) plates.
• Cathode: Lead dioxide (PbO2) coated plates.
• Electrolyte: Concentrated Sulfuric Acid (H2SO4, approx 38% by mass).

⇒Discharging Reactions (Using the battery)

// At Anode:
Pb + SO4^2- → PbSO4 + 2e^-

// At Cathode:
PbO2 + 4H^+ + SO4^2- + 2e^- → PbSO4 + 2H2O

// Notice that BOTH electrodes turn into Lead Sulfate (PbSO4)!
// Also, H2SO4 is consumed, and water is produced, lowering the acid density.

⇒Charging Reactions (Alternator running)

When plugged into a charger, the electrical current forces the reactions backwards. The PbSO4 on the anode turns back into spongy Pb. The PbSO4 on the cathode turns back into PbO2. H2SO4 is regenerated.

Lithium-ion batteries power the modern world (phones, laptops, EVs). Lithium is the lightest metal and has the highest oxidation potential, making it the ultimate material for high-energy-density batteries.

⇒The Rocking Chair Mechanism

Unlike lead-acid, Li-ion batteries do not rely on heavy chemical transformations. Instead, Lithium ions (Li+) physically 'rock' back and forth between the two electrodes through the electrolyte during charge and discharge. This process is called Intercalation.

• Anode: Porous Graphite (Carbon) holding intercalated Lithium atoms (LiC6).
• Cathode: A Lithium metal oxide lattice (like LiCoO2).
• Electrolyte: Lithium salts (LiPF6) dissolved in an organic solvent (water cannot be used, as Li reacts explosively with it).

⇒Discharge Process

// At Anode: Lithium atoms in graphite give up electrons and become Li+ ions
LiC6 → xLi^+ + xe^- + C6

// The Li+ ions swim through the electrolyte to the Cathode.
// The electrons travel through your phone circuit, powering it.

// At Cathode: The Cobalt oxide absorbs the arriving electrons and Li+ ions
Li(1-x)CoO2 + xLi^+ + xe^- → LiCoO2

A fuel cell is a galvanic cell that converts the chemical energy of a combustible fuel directly into electricity, without burning it. It operates continuously as long as fuel is supplied.

⇒Hydrogen-Oxygen Fuel Cell

Used heavily in the Apollo space missions. It takes Hydrogen gas as fuel and Oxygen gas as the oxidizer.

• Anode: Porous carbon infused with Platinum catalyst. H2 gas is pumped in.
• Cathode: Porous carbon infused with Platinum catalyst. O2 gas is pumped in.
• Electrolyte: Hot aqueous KOH solution.

// Anode: Hydrogen is oxidized
2H2 + 4OH^- → 4H2O + 4e^-

// Cathode: Oxygen is reduced
O2 + 2H2O + 4e^- → 4OH^-

// Overall Reaction: Pure water generation
2H2 + O2 → 2H2O + Electricity

Corrosion is the gradual, spontaneous, and destructive degradation of a metal due to chemical or electrochemical interaction with its environment.

⇒Why do metals corrode?

Metals exist in nature as stable ores (oxides, sulfides). We use immense thermodynamic energy in metallurgy to extract the pure metal. The pure metal is in an unstable, high-energy state. Corrosion is simply the metal returning to its stable, low-energy natural ore state.

⇒Dry (Chemical) Corrosion

Direct chemical attack by gases in the atmosphere (O2, Halogens, H2S) in the absolute absence of moisture.

When oxygen attacks, it forms a metal oxide layer. The nature of this layer determines the fate of the metal:

• Stable Layer (e.g., Al, Cr): Forms a tight, non-porous protective film. Corrosion stops immediately. This is why Aluminum doesn't rust away.
• Unstable Layer (e.g., Ag, Au): The oxide breaks down back into metal and oxygen. No corrosion occurs.
• Volatile Layer (e.g., Mo): The oxide layer evaporates as soon as it forms, exposing fresh metal. Rapid destruction.
• Porous Layer (e.g., Fe): The oxide (rust) has pores and cracks. Oxygen continues to penetrate, destroying the entire metal block.

Wet corrosion is vastly more common and destructive. It occurs when a conducting liquid (moisture) is in contact with the metal, or when two dissimilar metals are submerged in a solution.

⇒The Mechanism

Wet corrosion creates microscopic Galvanic Cells on the surface of the metal itself. One part of the metal acts as the Anode, another part as the Cathode, and moisture acts as the electrolyte.

• At the Anodic Area: The metal literally dissolves. Oxidation occurs, metal loses electrons and goes into the moisture as ions. (Destruction always happens at the anode).
• At the Cathodic Area: The electrons travel here through the metal. Reduction occurs. Oxygen from the air combined with water consumes the electrons to form Hydroxide ions.

// Anode (Iron dissolves):
Fe → Fe^2+ + 2e^-

// Cathode (Oxygen reduction):
½O2 + H2O + 2e^- → 2OH^-

// The Fe^2+ and OH^- meet in the moisture to form Rust:
Fe^2+ + 2OH^- → Fe(OH)2
Fe(OH)2 + O2 + H2O → Fe2O3·xH2O (Hydrated Ferric Oxide = RUST)

Electrochemical corrosion manifests in several distinct, dangerous ways depending on the environment.

⇒Galvanic Corrosion

Occurs when two DIFFERENT metals are physically connected and exposed to an electrolyte. The metal that is higher in the electrochemical series (more active/negative potential) automatically becomes the Anode and gets destroyed, while the other becomes the Cathode and is protected.

Example: If an Iron pipe is connected to a Copper pipe, Iron (more active) acts as the anode and corrodes rapidly. Copper is protected.

⇒Differential Aeration (Waterline) Corrosion

Occurs when a SINGLE piece of metal is exposed to different concentrations of Oxygen. The part of the metal exposed to LESS oxygen becomes the Anode. The part exposed to MORE oxygen becomes the Cathode.

Example: A steel post driven into the soil. The part deep in the soil has no oxygen (Anode = Corrodes). The part above ground has plenty of oxygen (Cathode = Protected). It breaks right at the soil line!

⇒Pitting Corrosion

A highly localized, extremely dangerous form of corrosion. A tiny scratch in a protective coating exposes a microscopic pinhole of metal. This tiny dot becomes the anode, while the massive coated surface becomes the cathode. The entire corrosive energy is focused on that tiny dot, drilling a deep hole straight through the metal.

Since corrosion causes billions of dollars in infrastructure damage annually, industrial engineers employ massive electrochemical defenses.

⇒Cathodic Protection

Since destruction ONLY occurs at the Anode, the fundamental strategy of protection is to force the structural metal to act as the Cathode.

⇒Metallic Coatings

• Galvanizing: Dipping Iron into molten Zinc. Zinc is more active. Even if the coating is scratched, the Zinc acts as a sacrificial anode, protecting the exposed Iron.
• Tinning: Coating Iron with Tin (used in food cans). Tin is LESS active than Iron. If the coating is perfectly intact, it protects. But if scratched, the exposed Iron becomes the anode and corrodes insanely fast.

⇒End-of-Unit Verification

• Distinguish clearly between Galvanic cells and Electrolytic cells (including Anode/Cathode polarities).
• Write the half-cell reactions and overall reaction for the Daniel Cell.
• Write the Nernst Equation and use it to calculate the EMF of a cell under non-standard concentrations.
• Explain the construction and working of the Standard Hydrogen Electrode (SHE) and state its drawbacks.
• Detail the charging and discharging chemical reactions of a Lead-Acid battery.
• Explain the 'Rocking Chair' mechanism of Lithium-ion batteries.
• Write the anodic and cathodic reactions for a Hydrogen-Oxygen Fuel Cell.
• Explain the electrochemical mechanism of wet corrosion (Rusting of iron) with equations.
• Describe Galvanic Corrosion and Differential Aeration Corrosion with real-world examples.
• Explain how a Sacrificial Anode provides cathodic protection to an underground pipeline.
• Compare the protective behavior of Galvanizing vs Tinning when the metal coating is accidentally scratched.